Acids and Bases

1. Acids and Bases Three major ways to define acids and bases introduced by Lewis, Brønsted and Arrhenius. They differ in the role of water Arrhenius and Brønsted require water, Lewis does not Acid Donates an H+ Brønsted Accepts an H+ Base Acid HCl + H2O → H3O+ + Cl- Base NaOH + H+→ Na+ + H2O

2. Acids and Bases Acid Produces H3O+ when added to water Arrhenius Produces OH- when added to water Base HCl + H2O → H3O+ + Cl- NH3+ H2O → NH4+ + OH-

3. : H H H H H B H H B N N H H H H H Acids and Bases Acid Accepts electrons Lewis Note: Electrons are not transferred between acids and bases, they are shared. Base Donates electrons BH3 + NH3 →BH3NH3 Base Acid

4. Acids and Bases The Lewis definition is the most general H+ leaves electrons behind Consider a Brønsted Acid It donates a H+ A-H i.e.A accepts the electrons A : _ H+ A is a Lewis Acid All Brønsted acids are Lewis Acids An Arrhenius acid, is a Bronsted Acid, since it produces H3O+ when dissolved in water as it “donates” H+ to H2O.

5. Strong Acids and Bases A strong acid, just like a strong electrolyte, is an acid which dissociates completely when dissolved in water. The concentration of H3O+ is thereby the highest possible, determined exactly by how much acid was added to water Ex) HCl + H2O → H3O+ + Cl- Ex) H2SO4 + H2O → H3O+ + HSO4_ Inorganic acids tend to be strong acids (except HF) A strong base, just like a strong electrolyte, is dissociates completely when dissolved in water. The concentration of OH_ is thereby the highest possible, determined exactly by how much acid was added to water. Ex) NaOH → Na+ + OH_ The hydroxides of alkali metals are strong bases.

6. Reactivity A reaction between an acid and a base produces water and a salt HCl(aq) + NaOH(aq)→ H2O (l) + NaCl (aq) Acid Base Water Salt Ionic Equation H3O+(aq) + Cl-(aq) + Na+(aq) + OH_(aq) → H2O(l) + Na+(aq) + Cl-(aq) Net equation H3O+(aq) + OH_(aq)→ H2O(l) A strong acid will react completely with any base. A strong base will react completely with any acid.

7. Weak Acids and Bases A weak acid, just like a weak electrolyte, does not dissociate completely when dissolved in water The concentration of H3O+ is not the highest possible, since much remains in the undissociated form. The concentration of H3O+ is determined from the dissociation constant, similar to Ksp, and the amount of acid added. Ex) CH3COOH (l) + H2O(l) → CH3COO-(aq) + H3O+(aq) Organic acid tend to be weak acids A weak base, just like a weak electrolyte, does not dissociate completely when dissolved in water. The concentration of OH- is not the highest possible, since much remains in the undissociated form. The concentration of OH- is determined from the dissociation constant, similar to Ksp, and the amount of acid added. Ex) NH3+ H2O → NH4+ + OH- Metal Oxides and nitrogen containing organic compounds tend to be weak bases

9. The oxides are anhydrides Non-metal oxides react with water to give oxoacids Therefore non-metal oxides are anhydrides of acids Metal oxides react with water to give hydroxide bases Therefore metal oxides are anhydrides of bases Metalloid oxides are amphoteric: they react with either strong acids or strong bases Amphotericoxides usually do not dissolve with water by themselves, but react with both strong acids and strong bases to give soluble products

10. Metal Metalliod Non-metal 1 2 13 15 16 14 17 The oxides are anhydrides Strength of acids and bases is correlated with the positions of the oxides on the PT

11. pH The acidity (or basicity) of a solution is reported as pH: pH = -log [H3O+] or [H3O+] = 10-pH = concentration of H3O+ in mol./l = molar (M) p = power of For pH < 7 solution is acidic For pH > 7 solution is basic Ex) 0.10 M solution of HCl [H3O+] = 0.10 M pH = - log [0.10] = -(-1.00) = 1.00 # of sig. figs. Increased from 2 to 3? For logarithmic quantities only the decimal numbers are significant. Therefore a pH = 1.00 has only 2 sig. figs, Note: pH does not have units

12. pOH Basicity of a solution can be reported as pOH: pOH = -log [OH-] Where [OH-] = conc. of OH in mol/l pH and pOH are related by: pH + pOH = 14 at 25 oC Therefore an acidic solution as pOH > 7, and a basic solution has pOH < 7 Determine the pH and pOH of: Exercise a) 0.275 M HNO3 solution [H3O+] = 0.275 M pH = - log (0.275) = -(0.561) = 0.561 pOH = 14.000 – pH = 14.000 -0.561 = 13.439 b) 0.0051 M NaOH solution [OH-] = 0.0051 M pOH = - log (0.0051) = -(-2.29) = 2.29 pH = 14.000 – pOH = 14.000 -2.29 = 11.71

13. Strength of an Acid • The quantity we use to measure the strength of an acid is its pKa • (corresponds to how easily an acid gives up H+). Acids with low pKa values • are strong while acids with high pKa values are weak: H2O (pKa = 14) WEAK ACIDS HCO3-1 (pKa = 10.3) H2CO3 (pKa = 6.4) pKa CH3CO2H (pKa = 4.7) Citric acid (pKa = 3.1) HF (pKa = 3.1) H3PO4 (pKa = 2.15) H3O+ (pKa = 0) This includes all aqueous solutions of HCl, HBr, HI, HNO3, HClO4 and H2SO4 Concentrated HNO3 (pKa = -1) Concentrated H2SO4 (pKa = -3) STRONG ACIDS Concentrated HCl (pKa = -7)

14. Aqua Complexes as Acids A hydrated proton has a pKa of 0 defining the line between the strong and weak acids. How about aqua complexes of metals? What factors affect how easily the aqua complex gives up H+?

15. Aqua Complexes as Acids If we plot pKa versus z2/r for a variety of aqua complexes, we see that there is a correlation. If we only look at those metals with low electronegativity values (1.5), we can approximate: If we introduce an empirical “fudge factor”, we get a more accurate – if more complex formula: Z2 r